This document discusses different types of batteries and their components and reactions. It provides details on primary batteries like Leclanche cell and mercury cell. It also describes secondary batteries like lead-acid battery and nickel-cadmium battery. Fuel cells and their working are explained. Corrosion and its prevention are discussed. Hydrogen economy and methods of hydrogen production are mentioned.

1. APPLICATION OF GALVANIC CELL: BATTERY
A battery is an electrochemical cell that is used as an energy sources. It stores energy in form of oxidising and
reducing agent and then releases energy electrically when it is needed. A practical battery should have following
characteristics:
 A practical battery should be reasonably light in weight and compact in size.
 A practical battery should give a constant voltage.
 A practical battery should be economic, easily transportable and having a longer life.
TYPES OF BATTERY
Primary Cell: In this type of cell the reaction occurs only once and the battery is then dead and cannot be used again.
i.e. in primary cells reaction occurs only in a single direction.
Secondary cell: It is type of cell which can be recharged by passing a current through them by a generator system and
they can be used again and again for charging and discharging.
PRIMARY BATTERY
LECLANCHE CELL: A dry cell is compact form of Leclanche cell. The cell consists of Zinc container that acts as
an anode. The zinc container is lined from inside with a porous insulating paper. The cathode is a carbon rod having a
brass cap . The space between cathode and anode is filled with a powdered mixture of MnO2 and a thick paste of
ammonium chloride, zinc chloride and charcoal powder. The porous paper lining prevents a direct contact between
Zinc container and paste, thus acts as a salt bridge
Anode (Oxidation): Zn (s) → Zn2+
+ 2e-
; zinc ions migrate towards carbon electrode.
Cathode (Reduction): 2MnO2(s) + 2NH4
+
+ 2e-
→ 2MnO(OH) +2NH3
Overall reaction: Zn(s) + 4MnO2(s) + 4 NH4
+
→ Zn2+
+ 2Mn2O3 + 4NH3 + 2H2O
Here MnO2 acts as a depolariser. A dry cell has a potential of about 1.5 V. They so called dry cells are not really dry.
These cells consists wet paste of NH4Cl and ZnCl2. A dry cell operate as long as the paste in it is moist.
Q. Why leclanche cell cannot be charged?
Ans: Due to following reaction
4 NH3 + Zn2+
→ [Zn(NH3)4]2+
Zn(s) + 2 NH4
+
→ Zn2+
+ 4 NH3 ↑ + H2 ↑
It is an acidic battery. Zn can be gradually eaten away of acidic NH4
+
.
MERCURY CELL
This cell provides a more stable voltage of emf 1.35 V. They are costlier and hence used in sophisticated devices ( low
current devices ) like camera, hearing aids, watches etc.
A mercury cell consists of Zn and Hg amalgam as anode which is covered by a steel top while a paste of HgO and
carbon powder as cathode. It is placed in contact with outré steel case.
Electrolyte: A paste of KOH and Zn(OH)2. This paste is carried in inter porous material.
Anode: Zn(Hg) + 2 OH-
→ ZnO(s) + H2O + 2e-
Cathode: HgO(s) + H2O +2e-
→Hg(l) + 2OH-
Overall reaction: Zn(Hg) + HgO(s) → ZN(OH)2 + Hg(l)
SECONDARY BATTERY: The secondary cells are repeated action cells. They can be recharged by passing
electricity through them. In these cells electrical energy is stored in form of chemical energy. Therefore they are also
called storage cells or accumulators.
LEAD STORAGE BATTERY
The batteries used in automobiles are lead storage batteries. Commercially in lead storage battery there are six or 12
lead storage cell connected together.
Each cell has anode and cathode plates separated from each other by separator. Separators are sheets of insulating
material placed in between plates. The plates in each electrode are called grids which are made up of lead or alloy of
lead and antimony. It consists of a lead anode and a grid of lead packed with lead dioxide as cathode. A 38% solution
of H2SO4 is used as an electrolyte.

2. All cathode plates (grids) marked as positive are coated with red-brown lead dioxide.
All anode plates (grids) marked as negative are coated with spongy lead.
Cell reaction:
Anode: PbS (s) + SO4
2-
→ PbSO4(s) + 2e-
Cathode: PbO2 (s) + SO4
2-
(aq) + 4H+
+ 2e-
→ PbSO4 (s) + 2H2O(l)
Overall reaction: PbS (s) + PbO2 (s) + 2H2SO4 (aq) → 2PbSO4 (s) + 2H2O(l)
The lead storage cell gives emf of 2 volts. So lead acid battery consisting of 6 cells will give 12 volt.
During use both the electrodes of lead storage cell get coated with white precipitate of PbSO4 and the H2SO4 gets
diluted by the water produced in the reaction. When both the electrodes het covered with lead sulphate, reaction stops,
and cell is said to be discharged or dead.
A discharged storage cell is recharged by passing a direct current source through storage cell. This direct current
reverses the overall reaction. During charging, the negative electrode of the storage cell is connected to the negative
side of dc source, and the positive electrode is connected to positive side of dc source.
When we charge our battery the cell reaction gets reversed and our overall reaction becomes:
charging
4 2 2 4
discharging
22PbSO (s) + 2H O(l) PbS (s) + PbO (s) + 2H SO (aq)
Ni-Cadmium Battery: (Alkaline Battery)
The nickel cadmium cell is also called nickel-cadmium accumulator/ the cell is represented as:
Cd(s) Cd(OH)2 Ni(OH)2 (s) NiO(OH)(s)
Electrolyte: aqueous solution of KOH.
Anode reaction: Cd(s) + 2OH-
↔ Cd(OH)2 +2e-
Cathode Reaction: 2[NiO(OH) + H2O +e-
→Ni(OH)2 + OH-
]
Overall Reaction: Cd(s) + NiO(OH) + H2O ↔ Cd(OH)2 + 2Ni(OH)2
The cell gives an emf of 1.35 V. Ni-Cd cells are expensive so are used high cost equipments.
Fuel Cell: A galvanic cell in which reactants are fed continuously to electrodes and products are removed
continuously from electrolyte compartment and it converts energy of combustion of fuels like H2, CH4, CH3OH etc
into electrical energy is called a Fuel Cell. Most common fuel cell is Hydrogen-Oxygen fuel cell.
Fuel Cell using H2 and O2
Catalyst used: Pt or Pd
Cell reaction: The reactions taking place in H2 - O2 fuel cell are
Anode : O2 + 4H+
+ 4e-
→ 2H2O E0
C = 1.229 V
Cathode: 2H2 (g) → 4H+
+ 4e-
E0
C = 0.00 V
Net Reaction: 2H2 (g) + O2 → 2H2O E0
C = 1.229 V
In this fuel cell, hydrogen and oxygen are bubbled through a porous carbon electrode into concentrated aqueous
sodium hydroxide solution. A fuel cell runs continuously as long as the supply of the reactants is continued.
Advantages of Fuel cell: Fuel cell like galvanic cells also produce electrical energy. Their advantage over galvanic cell
are:
1. Efficiency of fuel cell is 70% but under certain conditions their efficiency can be increased to 95%.
2. Fuel cell are pollution free i.e. why they are used in space crafts.
3. Fuel cell can supply energy continuously as long as fuels are fed into it without any drop in voltage.
Limitations of fuel cell:
1. High cost catalyst used in fuel cells.
2. Difficulties in handling gaseous fuels.
3. Highly corrosive nature of the electrolyte.

3. ELECTROCHEMICAL CORROSION:
It can be defined as spontaneous destruction of materials by chemical process. Electrochemical corrosion is a
spontaneous process where oxidation i.e. M→M+
+e-
is facilitated by presence of a suitable electron acceptor.
The rusting of iron, tarnishing of silver, development of green coating on copper and bronze are some examples of
corrosion.
The special characteristics of corrosion is that oxidation and reduction steps occur at separate location on the metals
because metals are conductive, so electrons can flow through the metal from anodic to cathodic region. The presence
of water acts a medium for transporting ions to and from the metal.
A corrosion system can be regarded as a short circuited electrochemical cell in which the
anodic process is like: Fe (s) → Fe2+
+ 2e-
cathodic process is like: O2+ 2H2O → 4e-
+ 4OH-
Corrosion: A two step process.
1. Fe (s) → Fe2+
+ 2e-
Fe atom at metal surface dissolves into moisture film leaving negative charge in the metal.
2. Corrosion continues as electron acceptor removes electrons from the metal. Common electron acceptors are:
O2+ 2H2O → 4e-
+ 4OH-
; Acid- 2H+
+ 2e-
→ H2 , Cu2+
+ 2e-
→ Cu
Factors Affecting corrosion:
1. More electropositive metals are more reactive. Therefore they get corroded faster. Fe gets corroded faster than
Ag because Fe is more electropositive than silver.
2. Pure metals do not corrode.
3. Reactive gases such as oxides of sulphur (SO2, SO3), oxides of nitrogen (NO2), and carbon (CO, CO2) in
presence of moisture produce acids and accelerate the rate of corrosion.
4. The presence of electrolytes like NaCl in water increases rate of corrosion. This is why Fe corrode faster in
sea water than in distilled water.
Prevention of corrosion:
1. By preventing surface of metallic object from atmosphere using paint of chemicals.
2. A metal can be prevented from corrosion by connecting it with a more electropositive metal. This method of
corrosion protection is called cathodic protection.
3. By covering metallic surface by other metals like Zn, Sn etc. which are inert to save the object. Tinning of
brass utensils is a very common practice in our country.
4. Metal surfaces are treated with phosphoric acid to form an insoluble phosphate coating on the surface.
Rusting Of Iron:
Rusting of iron is most common type of corrosion. Iron when exposed to moist air gets covered with a layer of brown
powdery material. Formation of brown powdery material on surface of iron in presence of moist air is called rusting
of iron. Rust mainly consists of hydrated ferric oxide (Fe2O3.xH2O)
The reaction which describes rusting of iron is: 4 Fe(s) + 3O2 + xH2O → Fe2O3.xH2O (s)
Rust is soft, porous, powdery substance. As a result of rusting iron loses its strength.
Hydrogen Economy
The energy demand of world is increasing day by day. The present source of energy i.e. fossil fuels are being depleted
fast. Efforts have been applied in developing unconventional and renewable sources of energy. Common renewable
source of energy are:
 Sunlight: Light energy of Sun is converted to electrical energy by Photovoltaic cell.
 Air: mechanical energy of wind is converted to electrical energy by wind mill.
 Tidal waves in sea: mechanical energy of wind is converted to electrical energy.
 Hydrogen: Hydrogen obtained from river or sea water can be used as fuel in a fuel cell.
J.O’M Bockris an American electrochemist with other scientists considered hydrogen as a versatile fuel for future
because its source i.e. water is in most abundance on earth and its combustion reaction to form water does not produce
pollution. Hydrogen can be obtained from water mainly by two methods:
1. Photoelectrolysis.

4. 2. Photochemical decomposition.
1. Calculate the emf of the cell in which the following reaction takes place:
Ni(s) +2Ag+
(0.002 M) → Ni2+
(0.160 M) + 2Ag(s). Given that E0
cell = 1.05V. Also calculate KC and ∆rG0
.
2. The cell in which the following reaction occurs:
2Fe3+
(aq) + 2I-
(aq) → 2Fe2
(aq) + I2 (s) has E0
cell = 0.236 V at 298 K. Calculate the standard Gibbs energy and
the equilibrium constant of the cell reaction. [AI 2012C]
3. The cell in which the following reaction occurs:
2Fe3+
(aq) 2I-
( aq)→ 2Fe2+
(aq) + I2(s) has E0
cell = 0.236 V at 298 K. Calculate the standard Gibbs energy and
the equilibrium constant of the cell reaction. [AI-2012C]
4. Under what condition is Ecell = 0, ∆rG = 0?
5. Express the relation among cell constant, resistance of the solution in the cell and conductivity of the solution.
How is molar conductivity of solution related to its conductivity? [AI=2012]
6. Why does the conductivity of a solution decrease with dilution?
7. Suggest a way to determine the value of
m


of water.
8. The molar conductivity of 0.025 molL-1
methanoic acid is 46.1 S cm2
mol-1
. Calculate its degree of dissociation
and dissociation constant. Given that λ∞
H+ = 349.6 S cm2
mol-1
and λ∞
CH3COO- = 54.6 S cm2
mol-1
.
9. The conductivity of a liquid held in a conductivity cell with electrodes of cross sectional area of 2 cm2
placed at a
distance of 1 cm from each other is 10-6
S cm-1
. Calculate cell constant and resistance of the conductivity cell.
[Answer: R = 5x105
ohm, κ = 10-6
S cm-1
]
10. The conductivity of a solution is 2.7 x 10-3
S cm-1
. This solution when filled in a conductivity cell offered a
resistance of 100 ohms. Calculate the cell constant of the conductivity cell. [Answer: 0.27 cm-1
]
11. The conductivity of 0.2 M solution of KCl at 298 K is 0.0248 S cm-1
. Calculate its molar conductivity. [Answer:
124 S cm2
/mol]
12. A 0.05 M NaOH solution offered a resistance of 31.6 ohm in a conductivity cell. If the cell constant of the
conductivity cell is 0.367 cm-1
then, determine the molar conductivity of the NaOH solution at this concentration.
[Answer: 232 S cm2
/ mol ]
13. The conductivity of 0.001 M acetic acid is 4 x 10-5
S cm-1
. Calculate the dissociation constant of an acid, if the
molar conductivity at infinite dilution for acetic acid is 390 S cm2
/mol. [Answer: α = 0.1026,Ka = 1.17x 10-5
]
14. The molar conductivity of 1.5 M solution of an electrolyte is found to be 138.9 S cm2
mol-1
. Calculate the
conductivity and resistivity of solution.
15. What products are obtained at cathode and anode when molten PbBr2 are electrolysed?
16. What are the products obtained during electrolysis of CuSO4 using Pt electrode?
17. How many faradays of charge are required to convert 1 mol of MnO4
-
to Mn2+
ion.
18. What mass of zinc will be reduced by 1 mole of electrons.
19. Write the chemistry of recharging the lead storage battery, highlighting all the materials that are involved during
recharging.
20. Explain how rusting of iron is envisaged as setting up of an electrochemical cell.
21. Calculate the emf of following cell:
Mg(s) Mg2+
(0.2 M) Ag+
(1x 10-3
M) Ag(s) . E0
Ag+/Ag = 0.80 V; E0
Mg2+/Mg= -2.37 V
22. The conductivity of 0.20 molL-1
solution of KCl is 2.48 x 10-2
S cm-1
. Calculate its molar conductivity and degree
of dissociation. Given λ0
(K+
)= 73.5 Scm2
mol-1
and λ0
(Cl-
) = 76.5 Scm2
mol-1
. [AI2015]
23. What type of battery is mercury cell? Why is it more advantageous than dry cell?
24. From the data of E0
values given below answer the following questions:
E0
M2+/M Cr Mn Fe Co Ni Cu
-0.91 -1.18 -0.44 -0.28 -0.25 +0.34
1. Why is E0
Cu2+/Cu value exceptionally positive?
2. Why is E0
Mn2+/Mn value highly negative as compared to other elements?
3. Which is a stronger reducing agent Cr3+
or Fe2+
? Why? [AI2015]
25. Name the type of cell which was used in Apollo space programme for providing electrical power. [AI2014]
26. What type of a battery is lead storage battery? Write the anode and cathode reactions and overall cell reaction
occurring in the operation of a lead storage battery. [AI2012C]
27. How many moles of mercury will be produced by electrolysing 1 M Hg(NO)3 solution with a current of 2.0 A for
3 hours. (M Hg(NO)3= 200.6 g/mol)